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This is essentially a question about the meaning and significance of the term vapour pressure (or vapor pressure if you're American). From what I understand a liquid in a container will have a certain amount of its vapour above it exerting a certain pressure: the vapour pressure - which is dependant on the identity of the substance and the temperature of the system only. The point at which the ambient pressure equals the vapour pressure the liquid boils - the liquid is in equilibrium with its vapour/gas phase (the line in the diagram separating liquid and gas).

However, this had me slightly confused because in the situation initially described (at a point somewhere within the liquid region - NOT on the phase boundary) some vapour existed above the liquid in the container and thus there must be some kind of equilibrium already existing. So, when the ambient pressure equals the vapour pressure nothing has changed in the sense that equilibrium remains present it's just that the pressure is no longer significant enough to facilitate the liquid.

Is what I've said above correct? Mainly in the definition of vapour pressure being the pressure of the vapour above the liquid AT ANY POINT in the liquid region on the phase diagram and also whether the vapour pressure is just a function of temperature.

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    $\begingroup$ Are you speaking of a situation in which there is a second component such as air present in the gas phase, or are you describing a situation in which there is only a single chemical component? $\endgroup$
    – Chet Miller
    Commented Sep 22, 2016 at 20:56
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    $\begingroup$ Does that make a difference? I think keep it simple and just consider a single component system and ideal gas behaviour. I'm not too concerned about the details of the system I just want to understand what vapour pressure really means $\endgroup$
    – RobChem
    Commented Sep 22, 2016 at 21:12
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    $\begingroup$ yes in a closed container vapour is at equilibrium with liquid (and vapour with solid if temperature/pressure low enough) at a given temperature. An open bottle of acetone, say, is not in equilibrium with the atmosphere, there just is not enough of it to become so. $\endgroup$
    – porphyrin
    Commented Sep 23, 2016 at 7:44
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    $\begingroup$ It does make a difference. It is easier to demonstrate that equilibrium does not necessary exist (and explain why) for a case in which there is a second component such as air in the gas phase. $\endgroup$
    – Chet Miller
    Commented Sep 23, 2016 at 12:13
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    $\begingroup$ Very few things in the real world are ever at equilibrium. But if an equilibrium has been established then the answer is yes. $\endgroup$
    – matt_black
    Commented Jul 3, 2017 at 23:51

3 Answers 3

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The definition of vapor pressure means that for any temperature and pressure within the limits of a liquid system, there is a certain amount of vapor present above the liquid. The liquid and vapor at this point are in equilibrium because the vapor pressure does not change over time for that given temperature and pressure. When the liquid starts boiling, the system is undergoing a phase transition and in not at equilibrium because the system volume is changing and energy is continually being added to raise the system's temperature. Of course, we know the temperature does not start to change until vaporization has completed.

Vapor pressure of any given pure liquid is normally a strong function of temperature and a weak function of pressure. If you took a beaker of a pure liquid and put it inside a container that could support a vacuum, then as the pressure external to the beaker starts to decrease, more molecules from the liquid phase have enough energy to escape the liquid phase and exist in the vapor phase. If you continued developing more vacuum, the liquid would boil. This is seen on the phase diagram as moving straight down from somewhere in the liquid region until you reach the gas-liquid equilibrium line.

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I know this question is very old, but the answers here are unsatisfactory, and so for the benefit for people with same questions in the future, let me clarify important points. I have assumed ideal scenario neglecting inter molecular forces.

Consider an open container with liquid phase of a substance with no molecule in gas phase above it, only air. At the top surface, there would be invariably some molecules with high energy due to their random motion. Therefore, they would begin to escape the liquid since nothing prevents them to do so. This is called evaporation. The liquid would continue to evaporate till any more evaporation would lead to partial pressure of its gas phase get in liquid region. This gas pressure is known as vapor pressure of the liquid, and is equal to maximum pressure in which the substance can remain in gas phase at that temperature. Note that liquid at this point has pressure equal to total gas pressure above it (atmospheric pressure).

Take a control volume inside the liquid. Remember that due to thermal motion, the molecules inside the control volume try to expand into gas phase, but since the gas phase pressure at that temperature is lower than the pressure by liquid, it can't expand. But as soon as the pressure exerted by the liquid becomes equal to the gas pressure, the molecules inside can easily transform into gas phase. (Note that it is not that straightforward--some seeding is required for bubble to form due to surface tension forces while the surrounding liquid is superheated, but I digress.)

And what is the pressure of the liquid equal to? Atmospheric pressure. But only when the container is open. So when container is open, and either atmospheric pressure or the temperature of the liquid is brought such that the pressure of the liquid becomes equal to its vapor pressure, then the thing boils.

If the container is closed, and all the air is removed, the pressure inside the liquid is equal to vapor pressure, and there will always be vapor pressure present, even if in minuscule amount. (You can't have a container filled with only liquid. Think about that for a moment. Similarly you can't have a container filled with only ice. Some water, or water vapour, or even combination of the two will appear alongside.) And in that case, any change in conditions of pressure or temperature will be immediately reflected by boiling or condensation of the system. That is, a closed container with only water inside can boil at any temperature.

So in summary, vapor pressure is the maximum gas pressure possible at that temperature. The liquid always tries to expand and attain vapor pressure, and in case the inside pressure is greater than the vapor pressure, the inside of the liquid can't do so, otherwise it would boil. However, the surface of the liquid, due to imbalanced forces, can do it at any point by evaporating, and thus forming gas above it till vapor pressure is attained.

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    $\begingroup$ Liquid (and solid) phases can only form if there are intermolecular forces. I assume you mean you're ignoring intermolecular forces just in the gas phase? $\endgroup$
    – hBy2Py
    Commented Jul 3, 2017 at 11:05
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    $\begingroup$ You can have a container filled with only liquid. You can have a container filled with only ice. Fill a zip lock bag under water, you have only liquid. Put it in the freezer, you have only ice after a while. It is best to use degassed pure water, otherwise there might be some bubbles of nitrogen etc. See the question this (in my mind) incorrect statement inspired: chemistry.stackexchange.com/questions/117850/… $\endgroup$
    – Karsten ♦
    Commented Jul 9, 2019 at 17:57
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You swirled around so much that it is hard to follow your logic.

The key here is the concept of equilibrium. If the system is at equilibrium then a liquid will have a particular vapor pressure at a particular temperature.

Think about it. The only reason that gasoline stays in the tank of a car (well at least old cars where the tank isn't sealed) is that the liquid gasoline is not in equilibrium with the atmosphere. If you just pour gasoline on the driveway and looked at the spot the next day then all of the gasoline will have evaporated.

So in general chemistry is about equilibrium conditions before and after some reaction. The major "exception" would of course be chemical kinetics.

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    $\begingroup$ May you please expand slightly? Is a liquid in equilibrium with it's vapour at some random value of P and T not on the phase boundary? $\endgroup$
    – RobChem
    Commented Sep 22, 2016 at 17:41
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    $\begingroup$ The point that I am trying to make is that you can have a system which has some pressure at some temperature which is not on the phase boundary because the system is not at equilibrium. "At equilibrium" is sometimes not true, like the gasoline in the tank of a car. $\endgroup$
    – MaxW
    Commented Sep 22, 2016 at 17:53
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    $\begingroup$ so a system can't exist in equilibrium at a point not on the phase boundary line? I don't really understand what you're saying $\endgroup$
    – RobChem
    Commented Sep 22, 2016 at 18:19
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    $\begingroup$ Let's ignore the solid phase for the moment. So we are just concerned with the liquid phase and the gas phase. At equilibrium there can be a gas phase with no liquid, but there cannot be a liquid phase with no gas phase (unless the container is absolutely full). So given the temperature, the curve for the gas-liquid phase boundary, the volume of the container, and the mass of the substance I can figure at equilibrium how much of the substance will be in the liquid form and how much will be in the gas form. Equilibrium is the key. $\endgroup$
    – MaxW
    Commented Sep 22, 2016 at 18:42
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    $\begingroup$ Look at your question - Is a liquid in a container always in equilibrium with its vapour? No, not always. It is very common to assume equilibrium but it isn't an absolute requirement that a system always be at equilibrium. // You can for example get pure water below below 0 Celsius in the liquid state because the system (liquid water & ice) is not at equilibrium. $\endgroup$
    – MaxW
    Commented Sep 22, 2016 at 18:58

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